Worked example 1
How many sigma and pi bonds are in a molecule of ethyne, C₂H₂?
Try it first: Draw the structure and mark every connection.
0 of 4 steps revealed.
BondingSigma and Pi bondsContent level: Core 16 min
Sigma and Pi Bonds
What you'll be able to do: Count sigma and pi bonds in a molecule and describe how each type of overlap forms.
Best after: Hybrid Orbitals from Domain Count
Introduction
Every bond you draw is one sigma bond plus zero, one or two pi bonds, and that split explains a lot of chemistry.
These are recommended, not required. You can start this lesson at any time.
Learning objectives
- Distinguish head-on from side-on orbital overlap
- Count sigma and pi bonds in a structural formula
- Relate pi bonds to unhybridized p orbitals
- Explain why sigma bonds are generally stronger than pi bonds
- Identify which orbitals form each bond in ethyne
Lesson
Sigma bonds
A sigma bond results from orbitals overlapping directly along the line joining the two nuclei. That head-on arrangement maximizes overlap, so sigma bonds are strong, and electron density is concentrated between the nuclei. Every bond between two atoms contains exactly one sigma bond, whether it is single, double or triple.
Count sigma bonds by counting lines between atoms once each, ignoring bond order.
Pi bonds
A pi bond forms when two parallel unhybridized p orbitals overlap sideways, creating electron density above and below the internuclear axis with a node along it. Side-on overlap is less effective than head-on overlap, so a pi bond is weaker than a sigma bond and is usually the first to break in a reaction. Only atoms that leave p orbitals unhybridized, meaning sp² or sp centres, can form pi bonds.
Counting in a real molecule
In ethyne, H-C triple bond C-H, each carbon is sp. There are two C-H sigma bonds, one C-C sigma bond and two C-C pi bonds, for three sigma and two pi bonds in total. In benzene, each of the six carbons is sp² with one p orbital, and those six p orbitals overlap around the ring in a delocalized pi system.
Delocalized pi systems spread electron density over several atoms and stabilize the molecule.
Key ideas
Definition
Sigma bond
A bond formed by head-on orbital overlap along the internuclear axis.
Definition
Pi bond
A bond formed by side-on overlap of parallel p orbitals, with a node along the internuclear axis.
Rule
Bond composition
Single is 1 sigma, double is 1 sigma and 1 pi, triple is 1 sigma and 2 pi.
Key concept
Ethene
C₂H₄ has 5 sigma bonds and 1 pi bond.
Equation
Sigma count
sigma bonds = number of bonded atom pairs
Equation
Pi count
pi bonds = total bond order - number of bonded atom pairs
Worked examples
Common mistakes
Counting a double bond as two pi bonds.
Why it's wrong: The first bond in any pair is always sigma.
Check instead: Subtract one from the bond order to get the pi count.
Claiming an sp³ atom can form a pi bond.
Why it's wrong: sp³ uses all three p orbitals, so none are left for side-on overlap.
Check instead: Check whether any unhybridized p orbital remains.
Saying pi bonds are stronger because double bonds are stronger.
Why it's wrong: A double bond is stronger than a single bond because it contains an extra bond, not because pi overlap beats sigma overlap.
Check instead: Compare pi to sigma directly: side-on overlap is less effective.
Practice this skill
No practice questions are available for this topic yet. You can still practice the whole unit.
What you should now know
A sigma bond forms by head-on overlap along the internuclear axis and is present in every bond. Pi bonds form by side-on overlap of unhybridized p orbitals above and below that axis. A single bond is one sigma, a double bond is one sigma and one pi, and a triple bond is one sigma and two pi bonds.
- Every bonded pair of atoms shares exactly one sigma bond
- Pi bonds come from side-on overlap of unhybridized p orbitals
- Single, double and triple bonds contain 0, 1 and 2 pi bonds
- Sigma overlap is more effective, so sigma bonds are stronger than pi bonds
- Only sp and sp² atoms have p orbitals free for pi bonding
Sources and further reading
This lesson is original Chem Help content. No external sources were adapted.